Coordination compounds: structure, bonding, and naming
Learn what coordination compounds are, how ligands bind to a metal center, how coordination numbers work, and how to name complexes step by step.
Quick answer
A coordination compound is a substance containing a central metal atom or ion bonded to surrounding molecules or ions, called ligands, that each donate at least one pair of electrons to the metal.
Example: Deep-blue [Cu(NH3)4]2+ forms when ammonia is added to a pale-blue copper(II) solution: four NH3 ligands bind the Cu2+ center, giving the tetraamminecopper(II) complex ion.
What you will learn
- Define a coordination compound and identify its central metal and ligands
- Determine the coordination number of a complex and predict its geometry
- Calculate the oxidation state of the central metal from the complex formula
- Name coordination compounds using standard nomenclature rules
- Distinguish the coordination sphere from counter ions in a formula
- Explain why chelating ligands form especially stable complexes
What a coordination compound is
A coordination compound is built around a central metal atom or ion with a group of molecules or ions, called ligands, attached directly to it. The metal together with its attached ligands forms a distinct structural unit, and that unit may be electrically neutral or carry a charge.
When the bracketed unit carries a charge, it is called a complex ion. The deep-blue species [Cu(NH3)4]2+, for example, is a complex cation: a copper(II) center surrounded by four ammonia ligands. Neutral complexes exist too, such as [Ni(CO)4], where a nickel(0) atom binds four carbon monoxide ligands.
Square brackets mark the coordination sphere: the central metal plus every ligand bonded directly to it. Anything written outside the brackets is a counter ion, not bonded to the metal.
The central metal and its ligands
Ligands act as Lewis bases: they donate an electron pair into empty orbitals on the metal, which acts as a Lewis acid. The bond that forms is a coordinate (dative) covalent bond, in which both shared electrons come from the ligand.
Transition metals are the usual centers because they offer vacant orbitals of suitable energy and can accept several ligands at once. The Swiss chemist Alfred Werner worked out this picture in the 1890s, showing that metals have a primary valence (oxidation state) and a secondary valence (the number of attached groups), an insight that earned him the Nobel Prize.
Coordination number and geometry
The coordination number is the number of donor atoms bonded directly to the central metal. It is counted by donor atoms, not by ligands: a bidentate ligand contributes two to the coordination number.
- Coordination number 2: linear geometry, as in [Ag(NH3)2]+
- Coordination number 4: tetrahedral (for example, [Zn(NH3)4]2+) or square planar (for example, [Pt(NH3)2Cl2])
- Coordination number 6: octahedral, the most common, as in [Fe(CN)6]4- and [Co(NH3)6]3+
The geometry matters enormously: it controls which isomers a complex can form and how its d orbitals split in energy, which in turn decides its color and magnetism.
Inside and outside the brackets
Formulas separate the coordination sphere from the rest of the compound. In [Co(NH3)6]Cl3, the six ammonia ligands are bonded to cobalt inside the brackets, while the three chloride ions sit outside as counter ions that balance the 3+ charge of the [Co(NH3)6]3+ complex.
Werner proved this distinction with conductivity and precipitation experiments: adding silver nitrate to [Co(NH3)6]Cl3 precipitates all three chlorides as AgCl, but the same test on [Co(NH3)5Cl]Cl2 precipitates only two, because the third chloride is bound inside the coordination sphere and does not behave as a free ion.
Finding the metal's oxidation state
The oxidation state of the metal is found by charge balance: the sum of the metal's oxidation state and the charges of the ligands inside the brackets must equal the charge of the complex ion.
Common neutral ligands include NH3, H2O, CO, and ethylenediamine. Common anionic ligands include Cl- (-1), CN- (-1), OH- (-1), and C2O4 2- (-2). In [Fe(CN)6]3-, six cyanide ligands contribute 6-, so iron must be +3 for the total to be 3-. In [Cu(NH3)4]2+, the ammonia ligands are neutral, so copper is +2.
Naming coordination compounds
Names are built systematically. Ligands are named first, in alphabetical order, with anionic ligands given an -o ending: chloro (Cl-), cyano (CN-), hydroxo (OH-), oxo (O2-). Neutral ligands keep special names: aqua (H2O), ammine (NH3, spelled with two m's), and carbonyl (CO).
The metal follows, with its oxidation state in Roman numerals. If the complex ion is an anion, the metal name takes an -ate ending: [Cu(NH3)4]SO4 is tetraamminecopper(II) sulfate, while K2[PtCl6] is potassium hexachloroplatinate(IV). Numerical prefixes (di-, tri-, tetra-, penta-, hexa-) count the ligands, and the number of complex ions in a salt is stated with a prefix too, as in tris(ethylenediamine)cobalt(III) chloride for [Co(en)3]Cl3.
Why coordination compounds matter
Coordination chemistry shows up everywhere once you know to look for it. The vivid colors of transition-metal complexes, their magnetic behavior, and their ability to catalyze reactions all trace back to the metal-ligand partnership.
The remaining articles in this series build on this foundation: ligands and the chelate effect, crystal field theory's explanation of color and spin states, the transition metals themselves, and the rich isomerism of complexes.
Variables and units
| Symbol | Name | Unit | Meaning |
|---|---|---|---|
| coordination number | dimensionless | Number of donor atoms directly bonded to the central metal center |
Common mistakes
The complex charge is the sum of the metal's oxidation state and all ligand charges. In [Fe(CN)6]4-, iron is in the +2 state even though the ion carries a 4- charge.
Brackets mark the coordination sphere. Species inside are bonded to the metal; species outside are counter ions with different chemical behavior, as Werner's precipitation experiments showed.
Many ligands are anions, such as Cl-, CN-, and OH-, and they contribute negative charge that must be included when finding the metal's oxidation state.
Neutral coordination compounds exist, for example [Ni(CO)4] and cis-[Pt(NH3)2Cl2], where the metal and ligand charges exactly cancel.
Real-world applications
Cisplatin, cis-[Pt(NH3)2Cl2], is a platinum complex used as an anticancer drug
EDTA chelation therapy treats heavy-metal poisoning by wrapping metal ions in a stable complex
Wilkinson's catalyst, a rhodium-phosphine complex, is used industrially for hydrogenation reactions
Prussian blue, an iron-cyanide coordination compound, is a classic deep-blue pigment
Gadolinium complexes serve as MRI contrast agents, improving image clarity in medical scans
Frequently asked questions
What is the difference between a complex ion and a coordination compound?
A complex ion is the charged bracketed unit itself, such as [Cu(NH3)4]2+. A coordination compound is the complete neutral substance, which may consist of the complex ion plus counter ions, as in [Cu(NH3)4]SO4, or may be a neutral complex with no counter ions at all, like [Ni(CO)4].
How do I find the oxidation state of the metal in a complex?
Add up the charges of the ligands inside the brackets, then find the metal oxidation state that makes the total equal the complex ion's charge. Remember that NH3, H2O, CO, and ethylenediamine are neutral, while Cl-, CN-, and OH- each contribute -1. In K4[Fe(CN)6], the six cyanides give 6- and the complex must be 4- to balance four K+, so iron is +2.
Why are the ligands written inside square brackets?
The brackets define the coordination sphere: everything inside is directly bonded to the central metal. Ions outside the brackets are counter ions. This distinction predicts real behavior, such as how many chloride ions precipitate with silver nitrate.
What are the most common coordination numbers?
Six is the most common, giving octahedral complexes such as [Co(NH3)6]3+. Four is next, giving tetrahedral complexes like [ZnCl4]2- or square planar ones like [Pt(NH3)2Cl2]. Two occurs in linear complexes such as [Ag(NH3)2]+. Higher numbers, such as 7 and 8, appear with larger metals and small ligands.
Why are so many coordination compounds brightly colored?
In transition-metal complexes, the ligand field splits the metal's d orbitals into different energy levels. Electrons can absorb visible light to jump between these levels, and the color you see is the complement of the light absorbed. Different ligands split the orbitals by different amounts, so changing ligands changes the color.